Why water boils at higher temperature at higher pressure

In summary, the higher pressure causes water to boil at a higher temperature because it resists the expansion of water molecules and requires a higher input of kinetic energy, which is heat. This is because pressure decreases the entropy of a gas, making it a less stable phase. The Clapeyron Equation explains this relationship between pressure and temperature. When pressure is increased, more energy is needed to overcome the outside pressure and turn the liquid into a gas.
  • #1
Tell me if this is correct:

the higher pressure causes water to boil at a higher temperature, therefore the bonds are harder to break because the atoms are experiencing resistance towards breaking free of the liquid and rising into gas.
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  • #2
not quite.

its the same reason why water evaporates at low pressure and why it condenses out of gas at high pressures. I am sure that you have seen high speed fighter jets and the compression vapor trails emanating from high pressure areas infront of the airframe or sonic boom compression waves...

anyway, as you heat something you are introducing kinetic energy at a molecular level. the kinetic energy is basically the vibrations of molecules, as they vibrate harder and push on their neighbors - they take up more volume and become transparent (aka gaseous), when something takes up more volume it increases displacement - when you increase displacement you increase buoyancy - that's why hotter gasses always float up... the same way hotter water floats above colder water...

however if you pressurize the boiling water container, the higher ambient pressure resists the expansion of the water molecules/vapour and thus requires a higher input of kinetic energy AKA heat.

boiling is nothing more than rapidly expanding water molecules jumping out of the bottom of your heated volume of water - propelled by buoyancy (because obviously the heat source is located on the bottom and heats/expands water molecules unevenly).
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  • #3
bobsmith76 said:
Tell me if this is correct:

the higher pressure causes water to boil at a higher temperature, therefore the bonds are harder to break because the atoms are experiencing resistance towards breaking free of the liquid and rising into gas.

No. The bonds take the same amount of energy to break.

One way to look at it is that boiling occurs because it increases the entropy of the universe (gases have higher entropy than liquids). But higher pressure decreases the entropy of a gas, because the atoms/molecules can assume fewer arrangements when they're packed more tightly. So the entropic gain is not as great when the pressure is increased, and the gas is no longer the stable phase. (In your example, liquid water would be more stable than water vapor at 102°C, for example.) Does this make sense?
  • #4
ok, I'm pretty sure I got it, here's my new explanation. I'm using this analogy. Imagine those pick three lotto machines from the 80s where the balls floated around in a cube. The balls are already moving and already have a lot of kinetic energy, so it would be easy to increase their movement with heat and thus turn them into the next phase, gas. Now imagine if you lowered the ceiling of the cube and slowly packed the balls into a tight space, then it would require more energy to break that ceiling and thus become gas.
  • #5
Seems like post 2 is a reasonable answer...but does not address the mechanism completely.

Post 4 does not make sense to me. Seems like a smaller volume would result in the balls hitting the ceiling more often increasing the likelihood of escape.

If you want to begin to understand boiling check here:

and here

If you think about volatility and how impurities affect "boiling" temperature, that is vapor pressure, you'll realize there is still more...salt water boils at a slightly higher temperature than fresh...

a good question for a chemistry expert.
  • #6
The equation which governs the slope of the liquid-vapour phase boundary is the Clapeyron Equation:
[tex]\frac{dp}{dT}=\frac{1}{T}(\frac{\Delta H}{\Delta V})_{Vaporisation}[/tex]

For vaporisation, [tex]\Delta H[/tex] is positive (you have to supply heat to vaporise a liquid; endothermic processes have positive [tex]\Delta H[/tex]) and [tex]\Delta V[/tex] is also positive (a gas occupies more volume than a liquid)

Therefore [tex]\frac{dp}{dT}[/tex] is positive; meaning that the temperature of boiling will increase with higher pressure.

Now, where the equation comes from is an argument involving entropy and chemical potential, but its a fairly standard concept and most good books on thermodynamics will have it. As an example, I have Atkin's Physical Chemistry, 8th edition.
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  • #7
consider the case of piston and cylinder assuming ideal conditions .a gas is taken in the cylinder provided with a piston (assuming intially equilibrium exist i.e pressure exerted by gas balances the weight of piston). now put more weight on the piston i.e you are exerting more pressure on the gas inside the cylinder ,due to this contraction will take place . now if you heat the cylinder i.e supply energy to the gas its kinetic energy will increase and so will the pressure exerted by the gas .heat it till piston reaches its orignal position. now let us analyse what has happened so far , placing weight on the piston increases the pressue exerted on the gas and to counter this increse in pressure energy was supplied to the gas . similarly during boiling when outside pressure is incresed we have to supply more energy to the liquid(water here) to increse the kinectic energy of its molecules to enough extent to over come the outside pressure.(during boiling water is converted into gassous phase)and we knoe boiling takes place when the vapour pressure of the liquid is equal to the pressure exerted by the outside environment.
  • #8
clearly your statement is wrong.

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